1. Bonding: General Concepts
Chapter 8

2. Types of Chemical Bonds-- Section 8.1
A chemical bond is formed by sharing or donating electrons between multiple atoms
Three main types of chemical bonds:
Ionic (No sharing of electrons)
Covalent (Electrons are shared albeit to different extents)
Metallic (Occurs only in purely metallic samples, e.g. iron, copper, gold)

3. Lewis Symbols
Valence electrons are typically used to form chemical bonds between atoms
Valence electrons of elements are represented using Lewis dot symbols

4. Writing Lewis Dot Symbols
Determine the number of valence electrons for the element in question
Arrange them around the element symbol, only pairing them after 4 have been placed
Examples:
Ca O Br Al

5. Ionic Bonding
Ionic bonding typically occurs between metals and nonmetals
Involves the complete transfer of an electron (no sharing)

6. Formation of Sodium Chloride

7. Typical Characteristics of Ionic Compounds
High melting points
Form crystalline (highly ordered) structures
Strong electrolytes
Brittle

8. Ionic Compounds and Lattice Energy
It is not really proper to speak of “molecules” of NaCl
Ionic molecules form an extended array of cations and anions that is referred to as a crystal lattice
Lattice energy is essentially the energy required to completely rip apart 1 mole of an ionic compound into gaseous ions
Coulombic forces for ionic crystal lattices play a major role:

9. Predicting Relative Lattice Energies
Predict which of the following would have the highest lattice energy out of each pair:
LiF or KF; LiCl or MgCl2
Explain why the lattice energy of Na2O is significantly higher than that of NaF.

10. Ion Formation and Electron Configuration
Why are ions formed in the first place?
Answer: Octet rule (obtaining noble gas electron configuration)
Ex: Reaction between Na and Cl
Na: 1s22s22p63s1 = [Ne]3s1
Na+ = [Ne]
Cl: 1s22s22p63s23p5 = [Ne]3s23p5
Cl-: 1s22s22p63s23p6 = [Ne]3s23p6 = [Ar]

11. Covalent Bonding Section 8.7
Covalent bonding occurs when the electrons forming a chemical bond are shared rather than transferred from element to another
While this is the main difference between ionic and covalent it is important to understand that there is essentially a spectrum between ionic and covalent compounds

12. Lewis Structures Section 8.10
Covalent bonds in covalent compounds are often represented by Lewis structures:

13. Multiple Bonds
A single bond is formed by sharing a single pair of electrons
Double bonds use two pairs and triple bonds three
As the bond order increases; the bond length decreases

14. Drawing Lewis Structures
Determine the number of electrons shared between all elements for a particular molecule (S = N – A)
Make all the necessary bonds using the shared electrons
Place lone pairs of electrons around atoms that do not have an octet of electrons
Exceptions H only needs 2; Be only needs 4; B only needs 6

15. Drawing Lewis Structures
Draw Lewis structures for the following molecules:
CCl4
CS2
OSCl2 (sulfur is central atom)
nitrate anion
See Interactive Example 8.6 (Pg. 328)

16. Bond Polarity and Electronegativity Section 8.3
The electrons that comprise a chemical bond between two elements are not necessarily shared equally
The degree to which electrons are shared is referred to as bond polarity
There are very few purely nonpolar covalent bonds
Ex: Any diatomic molecule (N2, O2, F2)

17. Trends in Electronegativity
The trend for electronegativity follows that of ionization energy:

18. Electronegativity and Bond Polarity
The difference in electronegativity between two elements determines the bond polarity
Although it is not a hard and fast rule, electronegativity differences greater than 2.5 usually indicate ionic compounds

19. Dipole Moments
As the difference in electronegativities increases, the bond becomes more and more polar and a dipole moment is created
Each bond in a compound could have a dipole moment
The net effect of each dipole moment adds up to an overall dipole for the entire molecule
Examples:
HF PFCl2 CCl4

20. Formal Charge
Formal charges for each atom must add up to the total charge on the molecule
Formal charge does not represent the actual charge on an atom, it is simply a method that is used to keep track of the electrons in a molecule
Example: Sulfate anion:

21. Lewis Structures and Formal Charge
When more than one Lewis structure is possible, the formal charges for each atom must be taken into account to provide the best Lewis structure:
The fewest number of formal charges represents the best Lewis structure
Negative formal charges should be placed on the atoms with the greatest electronegativities
Ex: thiocyanate ion (NCS-):

22. Resonance Structures Section 8.6
A molecule can have different resonance structures when two or more equal Lewis structures are possible
The bonds in resonance structures typically have lengths that are in between that of a single and double bond
Ex: nitrate ion (NO3-) and ozone (O3)

23. Sample Question
Write all possible resonance structures for the following species. Assign a formal charge to each atom. In each case which resonance structure is most important?
NO2- (N is central atom)
ClCN (C is central atom)

24. Exceptions to the Octet Rule Section 8.11
H, Be, and B have been previously mentioned as exceptions; however there are many examples of atoms with expanded octets
Elements capable of having expanded octets: S, I, Xe, P, Sb, Cl, As, Se, Br, Te
Ex: PCl5

25. How is Octet Expansion Possible?
The elements shown on the previous slide all have empty d-orbitals that are capable of accepting electrons for octet expansion
Ex: P

26. Drawing Lewis Structures w/ Expanded Octets
Draw Lewis structures for each of the species shown below indicating formal charge in all cases.
PF5
XeF2
SF6
See Interactive Example 8.7 (Pg. 331)

27. Expanded Octet Structures that Contain Oxygen
For compounds that have a central atom with an expanded octet and bonded to oxygen, you must see if increasing the bond order will result in a better Lewis structure
Examples: sulfate and sulfite ions

28. Odd Electron Species
There are a few molecules that have an odd number of valence electrons
As a result, it is impossible to place an octet of electrons around each atom
Ex: ClO2, NO, NO2

29. Strengths of Covalent Bonds Section 8.8
The bond enthalpy, H, is the enthalpy change associated with the breaking of a chemical bond
Example:
CH4(g) → C(g) + 4H(g) H = 1660 kJ

30. Average Bond Enthalpies for Common Covalent Bonds

31. Bond Enthalpies and Enthalpies of Reaction
The Hrxn can be estimated using Hess’s Law if the bond enthalpies for each chemical bond that’s broken or formed is known:
Hrxn = (Hbonds broken) - (Hbonds formed)

32. Using Average Bond Enthalpies
Using Table 8.5, estimate H for the following reaction:
CH4(g) + 2O2(g) → CO2(g) + 2H2O(g)
See Sample Exercise 8.5 (Pg. 324)

33. Molecular Structure: The VSEPR Model--Section 8.13
Lewis structures only provide a 2-D representation of a molecule
However, by including the bond angles of molecules, a more accurate 3-D representation can be achieved

34. VSEPR Theory
Valence Shell Electron-Pair Repulsion (VSEPR) theory states that the overall shape of a molecule is governed by the interactions of its electron clouds
The electron clouds for a given molecule repel each other and must separated as best as possible to minimize this repulsive interaction

35. The VSEPR Model
The electron domains included in the VSEPR model can be divided into two different categories:
Bonding pairs
Nonbonding (lone) pairs

36. Electronic vs. Molecular Geometries
The electronic geometry describes the shape of the electron clouds surrounding the central atom
Regardless of whether they are lone pairs or bonding pairs
The molecular geometry takes into consideration the effect of lone pairs but ignores their existence

37. Electronic and Molecular Geometry

38. Electronic and Molecular Geometry (cont.)

39. Determining the Shape of a Molecule
Draw Lewis structure
Determine number of lone pairs and bonding pairs (electronic geometry)
Remove lone pairs and determine molecular geometry

40. Examples
Draw Lewis structures and determine both electronic and molecular geometries for the following molecules:
H2O
ClNO (N as central atom)
CO32-
See Interactive Example 8.11 (Pg. 340)

41. Deviations from Ideal Bond Angles
The bond angles listed for the electronic geometries are ideal values
Several factors affect these ideal values
Existence of lone pairs
Large atoms
Bond orders >1
Lone pairs of electrons and bond orders >1 have the most significant impact

42. Geometries for Molecules w/ Expanded Octets
Trigonal pyramidal and octahedral geometries have two geometrically distinct points:
Equatorial
Axial

43. Examples
Draw Lewis structures and predict electronic and molecular geometries for each molecule shown:
BrF5
XeF4
ClF3
See Interactive Example 8.12 (Pg. 344)

44. Shapes of Larger Molecules
Molecular geometry can only be described with respect to a central atom
Molecules containing more than one central atom therefore have a different geometry about each particular atom

45. Bond Polarity (Revisited)
Predict whether the following molecules are polar or nonpolar:
CH4
HCN
H2O
XeF4

46. Covalent Bonding: Orbitals
Chapter 9
Covalent Bonding: Orbitals

47. Hybridization and the Localized Electron Model--Section 9.1
We know that Lewis structures provide a qualitative approach at determining molecular shape and that Schrodinger’s quantum #’s give us the shapes of atomic orbitals
These two concepts can be joined to form valence-bond theory

48. Orbital Overlap (VSEPR Theory)
Chemical bonds are simply represented with lines in a Lewis structure; however covalent bonds are formed from the overlap of atomic orbitals between two atoms

49. Bond Length

50. Hybrid Orbitals (Valence-Bond Theory)
Under valence-bond theory, the atomic orbitals used to create new chemical bonds mix together (hybridize) to create new hybrid orbitals
Example: BeF2

51. sp Hybridization (Two bonds)
The two new orbitals that are formed by mixing of the 2s and 2p orbitals are the hybrid orbitals
This allows Be to interact with two F atoms instead of one

52. Hybridization in BCl3 (Three bonds)
Hybridization of the 2s and two 2p orbitals leads to sp2 hybridization:

53. Hybridization in CCl4 (Four bonds)
Hybridization of the 2s and three 2p orbitals leads to sp3 hybridization:

54. Hybridization Involving d Orbitals
Elements capable of having an expanded octet use the empty d orbitals to do so; therefore hybrid orbitals can be created using d orbitals as well:
5 electron clouds  sp3d
6 electron clouds  sp3d2

55. Summary # of electron clouds Electronic Geometry Hybrid Orbital Set
Example 2
Linear sp
BeH2 3
Trig. Planar
sp2
BCl3 4
Tetrahedral
sp3
NH3 5
Trig. Bipyramidal
sp3d
PCl5 6
Octahedral
sp3d2
SF6

56. Examples
Write the Lewis structure, predict electronic and molecular geometries, and identify the hybrid orbital set for each molecule shown below.
SO3
NH4+
SF4
PF6-
See Interactive Example 9.5 (Pg. 364)

57. Multiple Bonds
All single bonds run along a line that passes through the nucleus of each atom and are known as sigma () bonds
A different bond forms when considering multiple bonds
Involves side to side overlap of two p orbitals to form pi () bonds

58. Hybridization in Ethylene and Acetylene